02 the chemical context of life


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  • Figure 2.3 The emergent properties of a compound.
  • Table 2.1 Elements in the Human Body
  • Figure 2.5 Simplified models of a helium (He) atom.
  • Figure 2.8 Energy levels of an atom ’s electrons.
  • Figure 2.9 Electron distribution diagrams for the first 18 elements in the periodic table.
  • Figure 2.10 Electron orbitals.
  • Figure 2.11 Formation of a covalent bond.
  • Figure 2.12 Covalent bonding in four molecules.
  • Figure 2.13 Polar covalent bonds in a water molecule.
  • Figure 2.14 Electron transfer and ionic bonding.
  • Figure 2.15 A sodium chloride (NaCl) crystal.
  • Figure 2.16 A hydrogen bond.
  • Figure 2.UN01 In-text figure, p. 41
  • Figure 2.17 Molecular shapes due to hybrid orbitals.
  • Figure 2.18 A molecular mimic.
  • Figure 2. UN02 In-text figure, p. 42
  • Figure 2.19 Photosynthesis: a solar-powered rearrangement of matter.
  • 02 the chemical context of life

    1. 1. LECTURE PRESENTATIONS For CAMPBELL BIOLOGY, NINTH EDITION Jane B. Reece, Lisa A. Urry, Michael L. Cain, Steven A. Wasserman, Peter V. Minorsky, Robert B. JacksonChapter 2The Chemical Context of Life Lectures by Erin Barley Kathleen Fitzpatrick© 2011 Pearson Education, Inc.
    2. 2. Matter consists of chemical elements in pureform and in combinations called compounds • Organisms are composed of matter • Matter is anything that takes up space and has mass© 2011 Pearson Education, Inc.
    3. 3. Elements and Compounds • Matter is made up of elements • An element is a substance that cannot be broken down to other substances by chemical reactions • A compound is a substance consisting of two or more elements in a fixed ratio • A compound has characteristics different from those of its elements© 2011 Pearson Education, Inc.
    4. 4. Figure 2.3 Sodium Chlorine Sodium chloride
    5. 5. The Elements of Life • About 20–25% of the 92 elements are essential to life, but it varies among organisms • Humans need 25 elements while plants only need 17 elements • Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter • Most of the remaining 4% consists of calcium, phosphorus, potassium, and sulfur • Trace elements are those required by an organism in minute quantities© 2011 Pearson Education, Inc.
    6. 6. Table 2.1
    7. 7. An element’s propertiesdepend on the structure of its atoms • Each element consists of unique atoms • An atom is the smallest unit of matter that still retains the properties of an element© 2011 Pearson Education, Inc.
    8. 8. Subatomic Particles • Atoms are composed of subatomic particles • Relevant subatomic particles include – Neutrons (no electrical charge) – Protons (positive charge) – Electrons (negative charge) • Neutrons and protons form the atomic nucleus • Electrons form a cloud around the nucleus • Neutron mass and proton mass are almost identical and are measured in daltons© 2011 Pearson Education, Inc.
    9. 9. Figure 2.5 Cloud of negative Electrons charge (2 electrons) Nucleus (a) (b)
    10. 10. Atomic Number and Atomic Mass • Atoms of the various elements differ in number of subatomic particles • An element’s atomic number is the number of protons in its nucleus • An element’s mass number is the sum of protons plus neutrons in the nucleus • Atomic mass, the atom’s total mass, can be approximated by the mass number© 2011 Pearson Education, Inc.
    11. 11. Isotopes • All atoms of an element have the same number of protons but may differ in number of neutrons • Isotopes are two atoms of an element that differ in number of neutrons • Radioactive isotopes decay spontaneously, giving off particles and energy© 2011 Pearson Education, Inc.
    12. 12. • Some applications of radioactive isotopes in biological research are – Dating fossils – Tracing atoms through metabolic processes – Diagnosing medical disorders© 2011 Pearson Education, Inc.
    13. 13. The Energy Levels of Electrons • Energy is the capacity to cause change • Potential energy is the energy that matter has because of its location or structure • The electrons of an atom differ in their amounts of potential energy • An electron’s state of potential energy is called its energy level, or electron shell© 2011 Pearson Education, Inc.
    14. 14. Figure 2.8 (a) A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons. Third shell (highest energy level in this model) Second shell (higher Energy energy level) absorbed First shell (lowest energy level) Energy lost Atomic nucleus (b)
    15. 15. Electron Distribution and ChemicalProperties • The chemical behavior of an atom is determined by the distribution of electrons in electron shells • The periodic table of the elements shows the electron distribution for each element© 2011 Pearson Education, Inc.
    16. 16. Figure 2.9 Hydrogen 2 Atomic number Helium 1H He 2He Mass number 4.00 Element symbol First shell Electron distribution diagram Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon 3Li 4Be 5B 6C 7N 8O 9F 10Ne Second shell Sodium Magnesium Aluminum Silicon Phosphorus Sulfur Chlorine Argon 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar Third shell
    17. 17. • Valence electrons are those in the outermost shell, or valence shell • The chemical behavior of an atom is mostly determined by the valence electrons • Elements with a full valence shell are chemically inert because it’s a very stable configuration© 2011 Pearson Education, Inc.
    18. 18. Electron Orbitals • An orbital is the three-dimensional space where an electron is found 90% of the time • Each electron shell consists of a specific number of orbitals© 2011 Pearson Education, Inc.
    19. 19. Figure 2.10 First shell Neon, with two filled Shells (10 electrons) Second shell (a) Electron distribution diagram First shell Second shell x y z 1s orbital 2s orbital Three 2p orbitals (b) Separate electron orbitals 1s, 2s, and 2p orbitals (c) Superimposed electron orbitals
    20. 20. The formation and function of moleculesdepend on chemical bonding between atoms • Atoms with incomplete valence shells can share or transfer valence electrons with certain other atoms • These interactions usually result in atoms staying close together, held by attractions called chemical bonds© 2011 Pearson Education, Inc.
    21. 21. Covalent Bonds • A covalent bond is the sharing of a pair of valence electrons by two atoms • In a covalent bond, the shared electrons count as part of each atom’s valence shell© 2011 Pearson Education, Inc.
    22. 22. Figure 2.11-3 Hydrogen atoms (2 H) Hydrogen molecule (H2)
    23. 23. • A molecule consists of two or more atoms held together by covalent bonds • A single covalent bond, or single bond, is the sharing of one pair of valence electrons • A double covalent bond, or double bond, is the sharing of two pairs of valence electrons • Covalent bonds can form between atoms of the same element of atoms of different elements© 2011 Pearson Education, Inc.
    24. 24. • The notation used to represent atoms and bonding is called a structural formula – For example, H—H • This can be abbreviated further with a molecular formula – For example, H2© 2011 Pearson Education, Inc.
    25. 25. Figure 2.12 Name and Electron Lewis Dot Space- Molecular Distribution Structure and Filling Formula Diagram Structural Model Formula (a) Hydrogen (H2) (b) Oxygen (O2) (c) Water (H2O) (d) Methane (CH4)
    26. 26. • Atoms in a molecule attract electrons to varying degrees • Electronegativity is an atom’s attraction for the electrons in a covalent bond • The more electronegative an atom, the more strongly it pulls shared electrons toward itself© 2011 Pearson Education, Inc.
    27. 27. • In a nonpolar covalent bond, the atoms share the electron equally • In a polar covalent bond, one atom is more electronegative, and the atoms do not share the electron equally • Unequal sharing of electrons causes a partial positive or negative charge for each atom or molecule© 2011 Pearson Education, Inc.
    28. 28. Figure 2.13 δ– O H H δ+ δ+ H2O
    29. 29. Ionic Bonds • Atoms sometimes strip electrons from their bonding partners • An example is the transfer of an electron from sodium to chlorine • After the transfer of an electron, both atoms have charges • A charged atom (or molecule) is called an ion© 2011 Pearson Education, Inc.
    30. 30. • A cation is a positively charged ion • An anion is a negatively charged ion • An ionic bond is an attraction between an anion and a cation© 2011 Pearson Education, Inc.
    31. 31. Figure 2.14-2 + – Na Cl Na+ Cl– Sodium atom Chlorine atom Sodium ion Chloride ion (a cation) (an anion) Sodium chloride (NaCl)
    32. 32. • Compounds formed by ionic bonds are called ionic compounds, or salts • Salts, such as sodium chloride (table salt), are often found in nature as crystals© 2011 Pearson Education, Inc.
    33. 33. Figure 2.15 Na+ Cl–
    34. 34. Weak Chemical Bonds • Most of the strongest bonds in organisms are covalent bonds that form a cell’s molecules • Weak chemical bonds, such as ionic bonds and hydrogen bonds, are also important • Weak chemical bonds reinforce shapes of large molecules and help molecules adhere to each other© 2011 Pearson Education, Inc.
    35. 35. Hydrogen Bonds • A hydrogen bond forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom • In living cells, the electronegative partners are usually oxygen or nitrogen atoms© 2011 Pearson Education, Inc.
    36. 36. Figure 2.16 δ– δ+ Water (H2O) δ+ Hydrogen bond δ– Ammonia (NH3) δ+ δ+ δ+
    37. 37. Van der Waals Interactions • If electrons are distributed asymmetrically in molecules or atoms, they can result in “hot spots” of positive or negative charge • Van der Waals interactions are attractions between molecules that are close together as a result of these charges • Collectively, such interactions can be strong, as between molecules of a gecko’s toe hairs and a wall surface© 2011 Pearson Education, Inc.
    38. 38. Figure 2.UN01
    39. 39. Molecular Shape and Function • A molecule’s shape is usually very important to its function • A molecule’s shape is determined by the positions of its atoms’ valence orbitals • In a covalent bond, the s and p orbitals may hybridize, creating specific molecular shapes© 2011 Pearson Education, Inc.
    40. 40. Figure 2.17 z Four hybrid orbitals s orbital Three p orbitals x y Tetrahedron (a) Hybridization of orbitals Space-Filling Ball-and-Stick Hybrid-Orbital Model Model Model (with ball-and-stick model superimposed) Unbonded Electron pair Water (H2O) Methane (CH4) (b) Molecular-shape models
    41. 41. • Biological molecules recognize and interact with each other with a specificity based on molecular shape • Molecules with similar shapes can have similar biological effects© 2011 Pearson Education, Inc.
    42. 42. Figure 2.18 Carbon Nitrogen Hydrogen Sulfur Natural endorphin Oxygen Morphine (a) Structures of endorphin and morphine Natural endorphin Morphine Endorphin Brain cell receptors (b) Binding to endorphin receptors
    43. 43. Chemical reactions make and breakchemical bonds • Chemical reactions are the making and breaking of chemical bonds • The starting molecules of a chemical reaction are called reactants • The final molecules of a chemical reaction are called products© 2011 Pearson Education, Inc.
    44. 44. Figure 2.UN02 2 H2 + O2 2 H2 O Reactants Reaction Products
    45. 45. • Photosynthesis is an important chemical reaction • Sunlight powers the conversion of carbon dioxide and water to glucose and oxygen 6 CO2 + 6 H20 → C6H12O6 + 6 O2© 2011 Pearson Education, Inc.
    46. 46. Figure 2.19
    47. 47. • Most chemical reactions are reversible: products of the forward reaction become reactants for the reverse reaction • Chemical equilibrium is reached when the forward and reverse reaction rates are equal© 2011 Pearson Education, Inc.
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